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What is the difference between real gas molecules and ideal gas molecules?
Kycesalis
Ideal gases are assuming that gas particles are discrete point particles, thus bouncing off each other with no attraction with one another, and each molecule taking up no space. This assumption allows for the Ideal gas law, which states exact proportions between measurable quantities in gases: pressure, volume, temperature, number of particles.
The ideal gas law is: PV = nRT
where:
P is pressure
V is volume
n is number of moles of gas
R is ideal gas constant
T is temperature (K)
Real gases particles, as common sense suggest, dohave volume and are minutely attracted to each other. Thus, gases do deviate from ideal behavior especially as they get more massive and voluminous. Thus, the attractions between the particles and the volume taken up by the particles must be taken into account. The equation derived by Van der Waals is the Van der Waals equation which simulates real gas behavior.
The Van der Waals equation is:
(p + ((n2a)/V2)(V - nb) = nRT
where:
p is measured pressure of the gas
n is number of moles of gas
a is attraction constant of the gas, varies from gas to gas
V is measured volume of the gas
b is volume constant of the gas, also varies from gas to gas
R is ideal gas constant
T is temperature (K)
Basically the Van der Waals equation is compensating for the non ideal attraction and volume of the gas. It is similar to PV = nRT, identical on the right side. To compensate for the massless volume that is found in ideal equation, the volume of the molecules are subtracted from the observed. Since, the equation of gas behavior concentrates on the space between the gas particles, and the volume of gas adds to the measured amount that should be used in the equation, thus it is subtracted from the equation. Another compensation is the fact that attraction between particles reduces the force on the walls of the container thus the pressure, thus it must be added back into the equation, thus the addition of the a term.
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